The Periodic table and it's periodicity

The periodic table is a systematic arrangement of elements based on their atomic structure and chemical properties. Its development was not the work of one person but a gradual process contributed to by many scientists.

The periodic table of elements is a common sight in classrooms, campus hallways and libraries, but it is more than a tabular organisation of pure substances. Scientists can use the table to analyse reactivity among elements, predict chemical reactions, understand trends in periodic properties among different elements and speculate on the properties of those yet to be discovered.

The first recorded attempt at creating a system to organise the elements was when Antoine Lavoisier published his table of elements in 1789. In 'Traite Elementaire de Chimie', Lavoisier listed 33 substances he considered elements, including light and caloric (heat).

Forty years later, in the year 1817, German physicist Johann Wolfgang Döbereiner observed similarities in physical and chemical properties of certain elements. He arranged them in groups of three in increasing order of atomic weight and called them triads, observing that some properties of the middle element, such as atomic weight and density, approximated the average value of these properties in the other two in each triad.

 

Law of Triads 

This law states that the atomic weight of the middle element is the arithmetic mean of the other two elements, which are arranged such that the elements having similar properties are placed in increasing order of atomic weight.

Examples:

 

Limitations

1. Döbereiner could identify only three triads as listed above from the elements known at that time.

2. The law does not apply to some elements in the same family, e.g F, Cl and Br.

 

Newland's Law of Octaves

In 1866, John Newlands established 56 known elements in increasing order of their atomic mass.

When elements are organised in the order of increasing atomic masses, the properties of the eight elements (starting from a given element) are a repetition of the properties of the first element, which is known as the law of octaves.

The octave's first and last notes are the same for the Indian music system: sa, re, ga, ma, pa, da, ni, sa. Similarly, in Newlands' table of octaves, the element F is eighth from the element H, implying that they have similar properties.

 

Limitations of Newland’s octaves

• It was only up until calcium that the classification of elements was completed through Newland’s Octaves.
• The invention of noble gases added to the constraints of this methodology since they couldn’t be included in this association without disturbing it fully.
• This arrangement had no provision for the future discoveries of elements.
• To adjust the existing order, Newlands has even placed two elements with different chemical and physical properties in the same position.

 

Mendeleeve Periodic Law

Dmitri Ivanovich Mendeléev, a Russian chemist, was the most important contributor to the early development of the periodic table. Many periodic tables were made, but the most important one was the Mendeleev periodic table. In 1869, following the rejection of Newlands' Law of Octaves, the Mendeleev Periodic Table emerged. In Mendeleev’s periodic table, elements were arranged based on the fundamental property, atomic mass, and chemical properties. During Mendeleev’s work, only 63 elements were known. After studying the properties of every element, Mendeleev found that the properties of elements were periodically related to atomic mass. He arranged the elements such that elements with similar properties fell into the same vertical columns of the periodic table. Later on, with the discovery of the electronic structure of the atoms, it became clear that elements do vary regularly not with their relative atomic mass, but with their atomic number.

 

 

Among chemical properties, Mendeleev treated formulae of hydrides and oxides as one of the basic criteria for categorisation. He took 63 cards and on each card, he wrote the properties of one element. He grouped the elements with similar properties and pinned them on the wall. He observed that elements were arranged in the increasing order of atomic mass, and there was the periodic occurrence of elements with similar properties. According to this observation, he formulated a periodic law which states: “the properties of elements are the periodic function of their atomic masses.”

In the Mendeleev periodic table, vertical columns in the periodic table and horizontal rows in the periodic table were named as groups and periods, respectively.

 

Merits of the Mendeleev Periodic Table

• Some gaps were left for the elements yet to be discovered. Thus, if a certain new element is discovered, it can be placed in a new group without disturbing any existing group.

Demerits of the Mendeleev Periodic Table

• He was unable to locate hydrogen in the periodic table.
• The increase in atomic mass was not regular while moving from one element to another. Hence, the number of elements yet to be discovered was not predictable.
• Later on, isotopes of elements were found that violated Mendeleev’s periodic law.

 

PLAY QUESTIONS

Q1. What were the advantages of Newlands’ law of octaves?

Answer:

• This law provided a framework for the classification of elements with similar properties in the same groups.
• This law was the first to be based on the atomic weight of elements, thereby linking atomic weight to the elemental features.
• This method performed well for lighter elements, for example, combining lithium, sodium and potassium in the same group.

Q2. Based on Newlands’ classification, the properties of sulphur are similar to those of oxygen because sulphur is the ______ element starting from oxygen.

 A. 7th
B. 8th
C. 3rd
D. 6th

Answer: Newlands’ law states similarity in the first and its corresponding 8th element. Hence, the answer is option B.

Q3. A and B are two elements having similar properties which obey the law of octaves. How many elements are there between A and B?

 A. 6
B. 5
C. 8
D. 7

Answer: Since the first and the eighth elements have similarities, there will be 6 elements between A and B. Hence, the answer is option A.

Q4. Which of the given pairs follows Newlands’ law of octaves?

 A. Boron and Sodium
B. Lithium and Sodium
C. Calcium and Potassium
D. Calcium and Nitrogen

Answer: Newlands’ law stated that every eighth element had similar properties when they were arranged in the increasing order of their atomic masses. Following this rule, lithium and sodium, sodium and potassium, chlorine and fluorine have similarities. So the answer is Option B.

Q5. John Newlands' Law of Octaves worked for all known elements at the time. How true is this?

Q6. ____________'s Law of Triads grouped elements into sets of three with similar chemical properties.
Q7. The Law of Octaves, which stated that every eighth element had similar properties, was proposed by ____________.

Q8. What key innovation did Mendeleev use in his periodic table?

A. He arranged elements by atomic number.

B. He left gaps for undiscovered elements.

C. He ignored element properties.

D. He included only noble gases.

 

 

THE MODERN PERIODIC TABLE AND MOSELEY'S PERIODIC LAW

In the year 1913, Henry Moseley studied the frequencies of the X-rays which were emitted when certain metals were bombarded with high-speed electrons. He found that in all the cases, the square root of the frequency was directly proportional to the atomic number of the atom of the metals. These studies believe that atomic number is the fundamental property of an element. In the above observation, Moseley gave the “Modern Periodic Law”, which states that: Physical and chemical properties of the elements are the periodic functions of their atomic numbers.

 

Features of The Periodic Table

The modern form of the periodic table is divided into eighteen vertical columns known as GROUPS, and seven horizontal rows known as PERIODS.

 

 

GROUP

A group is a vertical column of the periodic table, based on the organisation of the outer shell electrons. There are a total of 18 groups. Two different numbering systems are commonly used to designate groups, and you should be familiar with both. The traditional system used in the United States involves the use of the letters A and B. The first two groups are 1A and 2A, while the last six groups are 3A through 8A. The middle groups use B in their titles. Unfortunately, there was a slightly different system in place in Europe. To eliminate confusion, the International Union of Pure and Applied Chemistry (IUPAC) decided that the official system for numbering groups would be a simple 1 through 18 from left to right. Many periodic tables show both systems simultaneously.

 

PERIOD

A period is a horizontal row of the periodic table. There are seven periods in the periodic table, with each one beginning at the far left. A new period begins when a new principal energy level begins filling with electrons. Period 1 has only two elements (hydrogen and helium), while periods 2 and 3 have 8 elements. Periods 4 and 5 have 18 elements. Periods 6 and 7 have 32 elements, because the two bottom rows that are separate from the rest of the table belong to those periods. These two rows are pulled out in order to make the table itself fit more easily onto a single page.

 

PERIODIC CLASSIFICATION INTO BLOCKS AND FAMILIES

The elements in the periodic table may be divided into blocks according to the orbitals their valence electrons are found, which is responsible for the positions of the elements. The s-block elements have s-electrons in the outermost energy level, while the p-block has both s and p-electrons. The transition elements contain d-electrons in addition to their s and p-electrons, while the lanthanides and actinides contain f-electrons in addition to the s, p and d electrons.

 

 

FAMILIES OF ELEMENTS

Elements in the same group may be said to belong to a family since they show similar properties because their atoms have the same number of valence electrons. At the same time, certain properties of the element in the same group show a gradual change with an increase in atomic number. Such a gradual change of property within a group is known as GROUP TREND.

 

GROUP I (1)

The group I elements include: Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Caesium (Cs), and Francium (Fr). They are univalent elements. Their properties are as follows:

·         They are good reducing agents since they can readily donate one electron to form a cation.

·         They are metals, thus they are good conductors of electricity and heat.

·         They react vigorously with cold water to liberate hydrogen gas and form alkali; hence, they are known as ALKALI METALS. Example2Na(s) + 2H2O(l)→ 2NaOH(aq) + H2(g)

·         The oxides of group I elements dissolve in water to give very strong alkalis.

Example: K₂O(s) + H₂O(l) → 2KOH(aq)

 

GROUP II (2)

Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra) belong to group II. They are divalent elements and are also known as alkaline earth metals.

Their properties include:

·         They ionise by donating their two valence electrons; hence, they are good reducing agents.

·         They are hard metals, ductile and malleable and can conduct both electricity and heat.

·         Beryllium does not react with cold water or steam, magnesium reacts with steam only, while calcium reacts slowly with cold water to liberate hydrogen gas.

·         Their oxides are insoluble in water except for calcium oxide, which dissolves in water to form an alkali.

Example:  CaO(s) + 2H₂O(l) → Ca(OH)₂(aq)

 

GROUP III (13)

The group III elements are: Boron (B), Aluminium (Al), Gallium (Ga), Indium (In) and Thallium (Tl). They are trivalent elements.

Their properties are:

·         They are reducing in nature since they can donate their three electrons to form electrovalent compounds.

·         Only aluminium can react with steam at about 750 °C to liberate hydrogen gas.

·         Oxide and hydroxide of aluminium are amphoteric, i.e, they have both acidic and basic properties.

Examples:

Al₂O₃(s) + 3H₂SO₄(aq) → Al₂(SO₄)₂(aq) + 3H₂O(l)

2Al(OH)₃(s) + NaOH(aq) → NaAl(OH)₄(aq)

 

GROUP IV (14)

Group IV elements include: Carbon (C), Silicon (Si), Germanium (Ge), tin (Sn) and lead (Pb).

·         They form covalent compounds.

·         They exhibit two oxidation states: +2 and +4. Due to the inert pair effect of electrons in the s-orbital of the valence shell, the +2 oxidation state becomes more prominent down the group.

·         Electropositivity increases down the group. Carbon is a non-metal; silicon and germanium are metalloids, while tin and lead are metals.

·         Carbon does not react with water in any form, but silicon and tin react with steam at red heat to form +4 state oxides and hydrogen.

Example:  Si(s) + 2H₂O(l)→ SiO₂(s) + 2H₂(g)

GROUP V (15)

Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb) and Bismuth (Bi) belong to group V.  They have the following properties:

·         They exhibit oxidation states of -3 and -5.

·         They also show group trends. Nitrogen and phosphorus are non-metals; arsenic and antimony are metalloids, while bismuth is a metal.

·         They are electron acceptors, hence they are oxidising in nature.

·         They form oxides that dissolve in water to form acids, except nitrogen (I) oxide.

 

GROUP VI (16)

Elements in group VI include: Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te), and Polonium (Po). Their properties are as follows:

·         They are non-metals and exist as solids at room temperature except for oxygen.

·         They are electron acceptors and oxidising in nature.

·         They do not react with water in any form, but oxygen and sulphur combine directly with hydrogen to yield water and hydrogen sulphide, respectively.

 

GROUP VII (17)

Elements in this group include: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I) and Astatine (At). They are known as halogens (salt-makers). Their properties include:

·         They ionise to form univalent anions.

·         They exist as diatomic molecules.

·         As electron acceptors, all halogens are good oxidising agents.

·         They exhibit a group trend. Fluorine and chlorine are gases, bromine is a liquid and iodine and astatine are solids at room temperature.

 

GROUP VIII (0) or (18)

The elements in group 0 are known as rare or noble gases because they are non-reactive and exist freely as monoatomic molecules in the atmosphere. The elements that belong to this group are:

Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe) and Radon (Rn).

 

TRANSITION ELEMENTS

These are elements found in-between group II and III of the periodic table. The first transition series consists of elements: Scandium (Sc), Titanium (Ti), Vanadium (V), Chromium (Cr), Manganese (Mn), Iron (Fe), Cobalt (Co), Nickel (Ni), Copper (Cu) and Zinc (Zn).

Transition elements have the following properties:

1. High tensile strength

2. High melting and boiling points

3. Variable oxidation states

4. Formation of colored ions

5. Formation of complex ions

6. Paramagnetic in a mixture

7. Catalytic ability

 

LANTHANIDES (RARE EARTH ELEMENTS): These are found in period six. This series begins with Lanthanum (La) and ends with Lutetium (Lu). There are altogether 15, and they resemble one another greatly.

 

ACTINIDES AND THE ARTIFICIAL ELEMENTS: The actinides are similar to the Lanthanides. They are found in the seventh period, which starts with Actinium (Ac) and ends with Lawrencium (Lr). The famous Uranium is in this group.

The elements with atomic numbers from 93 to 103 are known as the artificial elements. This is because they do not occur naturally but were formed during nuclear reactions.

 

ELECTRONIC CONFIGURATION

The electronic configuration of an atom is the representation of the arrangement of the electrons distributed among the orbital shells and subshells. Commonly, the electronic configuration is used to describe the orbitals of an atom in its ground state.

 

MEANING OF ATOMIC ORBITAL

Orbital is the region of space around the nucleus where there is a high probability of finding an electron. The

four different types of orbitals are s, p, d, and f. These orbitals have different shapes, and one orbital can hold

a maximum of two electrons. The p-orbital has three degenerate orbitals, with a maximum of six electrons; the d-orbital has five sub-orbitals with a maximum of ten electrons, and the f-orbital has seven sub-orbitals with a maximum of fourteen electrons.

 

 

RULES AND PRINCIPLES FOR FILLING ELECTRONS IN THE ORBITALS

1. Aufbau Principle

According to the Aufbau Principle, electrons are added to atomic orbitals in order of increasing energy levels. This means that electrons will first occupy the lowest energy orbital available before moving to higher energy orbitals. The common order of orbital filling is:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

2. Pauli Exclusion Principle

The Pauli Exclusion Principle states that no two electrons in the same atom can have the same set of four quantum numbers. As a result, an orbital can hold a maximum of two electrons, and these two electrons must have opposite spins—one with an upward spin (↑) and one with a downward spin (↓).

3. Hund’s Rule

Hund’s Rule of Maximum Multiplicity explains how electrons are distributed within orbitals of the same energy level, such as the three orbitals in the p-sublevel or the five in the d-sublevel. It states that electrons will occupy each degenerate (equal-energy) orbital singly before any orbital gets a second electron.

Note: When writing electron configurations, these rules work together to ensure that electrons fill orbitals as stably and efficiently as possible.

The electronic configuration of some elements is given below for better understanding.

PERIODICITY (Periodic Trends)

Periodicity of elements refers to the repeating patterns in the chemical and physical properties of elements when they are arranged in order of increasing atomic number. This phenomenon is observed in the periodic table and is due to the recurring similar valence shell electronic configurations.

 

 

 

PROPERTIES OF ELEMENTS THAT SHOW PERIODIC VARIATION

1. ATOMIC RADIUS: This is the distance from the nucleus (centre) of an atom to the outermost orbital of its electron (shell).  The atomic radius is also defined as one-half the distance between two covalently bonded atoms. For electrovalent compounds, ionic radius is a measure of the distance between the centre of the ion and the centre of its nearest neighbour of opposite charge.

Trend

Atomic radius decreases across the periods from left to right, while it increases down the groups from top to bottom.

2. IONIZATION ENERGY: Ionization energy is the energy required to remove one mole of electrons from one mole of gaseous atom to produce one mole of gaseous ions.

Trend

Ionization energy increases across the periods left to right, while it decreases down the groups from top to bottom.

3. ELECTRONEGATIVITY: Electronegativity is the ability or tendency of an atom of an element to attract electrons to become negatively charged.

Trend

Electronegativity increases from left to right across each period, while it decreases down the groups from top to bottom.

4. ELECTROPOSITIVITY: Electropositivity is the power of the atom of an element to lose an electron and become positively charged.

Trend

Electropositivity decreases from left to right across the period, while it increases down the groups

from top to bottom.

5. ELECTRON AFFINITY: Electron affinity is the energy change which accompanies the addition of one mole of electrons to one mole of gaseous atom of an element to form negatively charged ions.

Trend

Electron affinity decreases from left across each period while it increases down the groups.

6. METALLIC CHARACTER: The metallic character of an element can be defined as how readily an atom can lose an electron.

Trend

Metallic property decreases across the periods and increases down groups, while non-metallic property increases across periods and decreases down groups.

7. CHEMICAL REACTIVITY: The tendency of an atom to undergo a chemical reaction.

Trend

 Chemical reactivity decreases from group I–III for metals and increases from group IV to VII for non-metals across the periods, while it increases down groups for metals and decreases down groups for non-metals.